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Oxidation and reduction are terms used to describe the movement of electrons in a compound/reaction mixture. Each atom, ion or molecule within a system is assigned an oxidation state, based on how many electrons it contains. If the species loses some of these electrons, it is said to have been oxidised. If it gains electrons it is said to have been reduced.
Remember OIL RIG
OIL: Oxidation Is Loss (of electrons)
RIG: Reduction Is Gain (of electrons)
A redox reaction involves two or more species being oxidised or reduced simultaneously, with the net result being that atom(s)/ ion(s) undergo a change in oxidation state. In these cases, the species being reduced is referred to as an oxidising agent, whilst the species being oxidised is referred to as a reducing agent. Reduction and oxidation cannot occur independently from each other but we can treat each part of the reaction independently by writing out half equations.
For example, in the formation of hydrogen chloride from hydrogen and chlorine:
H2 + Cl2 –> 2HCl
In this reaction, the starting materials are both neutral molecules and as a result, both have oxidation states of 0. However, when they react with each other, they form the ions H+ and Cl–. Each hydrogen atom has lost an electron, whilst each chlorine atom has gained one. Going back to OIL RIG, this must mean that hydrogen has been oxidised whilst chlorine has been reduced. Hydrogen now has an oxidation state of +1 and chlorine has an oxidation state of -1.
This is a bit clearer when you consider the half reactions:
H2 –> 2H+ + 2e– OXIDATION
Cl2 + 2e– –> 2Cl– REDUCTION
Putting these together gives:
H2 + Cl2 + 2e– –> 2H+ + 2Cl– + 2e–
Cancelling out the electrons gives:
H2 + Cl2 –> 2H+ + 2Cl–
Where the ions can react together to form 2HCl.
It’s not just in the lab that we see oxidation and reduction reactions. The most common oxidation we can observe is rust, which occurs through a series of redox reactions between iron, water and oxygen.
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